Question

(a) Draw the Lewis structures for at least four species that have the general formula $$ [: \mathrm{X} \equiv \mathrm{Y}:]^{n} $$ where \(X\) and Y may be the same or different, and \(n\) may have a value from \(+1\) to \(-2 .(\mathrm{b})\) Which of the compounds is likely to be the strongest Brønsted base? Explain.

Short answer

The four species with Lewis structures are: 1. \( \mathrm{N \equiv N} \): Both nitrogen atoms have a lone pair of electrons and no net charge. 2. \([:\mathrm{C} \equiv \mathrm{N}:]^{-} \): Carbon atom has one lone pair, nitrogen atom has two lone pairs, and the net charge is -1. 3. \([:\mathrm{N} \equiv \mathrm{O}:]^{+} \): Nitrogen atom has two lone pairs, oxygen atom has one lone pair, and the net charge is +1. 4. \([:\mathrm{O} \equiv \mathrm{O}:]^{2-} \): Both oxygen atoms have two lone pairs each and the net charge is -2. Among these species, \([:\mathrm{O} \equiv \mathrm{O}:]^{2-}\) is likely to be the strongest Brønsted base due to its high electron density from two negative charges, making it more favorable for proton acceptance.

Step-by-step solution

Understand the general formula

The general formula is given as: $$ [: \mathrm{X} \equiv \mathrm{Y}:]^{n} $$ Here, X and Y are atoms which may be the same or different, and n is the net charge on the species, which can range from +1 to -2. The formula shows a triple bond between X and Y, and lone pairs on both X and Y.

Draw the Lewis structures for four different species

We need to find four different species fitting the given formula and draw their Lewis structures. Let's consider the following combinations: 1. X = N, Y = N, n = 0 (Neutral molecule) 2. X = C, Y = N, n = -1 (Anion) 3. X = N, Y = O, n = +1 (Cation) 4. X = O, Y = O, n = -2 (Anion) Now let's draw the Lewis structures for these species: 1. \( \mathrm{N \equiv N} \): Both nitrogen atoms will have a lone pair of electrons and no net charge. 2. \([:\mathrm{C} \equiv \mathrm{N}:]^{-} \): Carbon atom will have one lone pair and nitrogen atom will have two lone pairs, and the net charge will be -1. 3. \([:\mathrm{N} \equiv \mathrm{O}:]^{+} \): Nitrogen atom will have two lone pairs and oxygen atom will have one lone pair, and the net charge will be +1. 4. \([:\mathrm{O} \equiv \mathrm{O}:]^{2-} \): Both oxygen atoms will have two lone pairs each and the net charge will be -2.

Determine the strongest Brønsted base

Brønsted bases accept protons (H+ ions). A stronger base will have a greater tendency to accept protons. In general, the higher the electron density on the atom that will bond with the proton (usually the atom with the lone pair), the stronger the base. Let's analyze the electron density of each species: 1. \( \mathrm{N \equiv N} \): Nitrogen atoms have no net charge, moderate electron density. 2. \([:\mathrm{C} \equiv \mathrm{N}:]^{-} \): Carbon atom has a negative charge, higher electron density. 3. \([:\mathrm{N} \equiv \mathrm{O}:]^{+} \): Nitrogen atom has a positive charge, lower electron density. 4. \([:\mathrm{O} \equiv \mathrm{O}:]^{2-} \): Oxygen atoms have a negative charge, high electron density. Comparing the electron density for each species, we can conclude: The compound \([:\mathrm{O} \equiv \mathrm{O}:]^{2-}\) is likely to be the strongest Brønsted base, as it has a high electron density due to the two negative charges, which will facilitate proton acceptance.

Key concepts

Brønsted Base

A Brønsted base is a substance that can accept protons (hydrogen ions, H+). In chemical reactions, Brønsted bases act as proton acceptors, often resulting in the formation of their conjugate acids. The ability of a Brønsted base to accept a proton depends on several factors, including the availability of a lone pair of electrons that can be used to bond with the proton.

When analyzing compounds to determine their strength as Brønsted bases, the key aspect to consider is the electron density around the atom that will bond with the hydrogen ion. Higher electron density makes it easier for the base to accept a proton, thus functioning as a stronger base. For instance, in species with high negative charges, the electron density is increased, which enhances their basicity.

In the exercise involving various species with the general formula \[[: \mathrm{X} \equiv \mathrm{Y}:]^{n}\], the strongest Brønsted base would be the species with the greatest negative charge, leading to the highest electron density. As we saw, out of the options provided, the \([:\mathrm{O} \equiv \mathrm{O}:]^{2-}\) ion, having two negative charges, would be the strongest Brønsted base, ready to accept a proton readily due to its significant electron-rich environment.

Electron Density

Electron density refers to the probability of finding electrons in a certain area around the nucleus of an atom or across a bond. This concept is pivotal in determining the chemical behavior of atoms within a molecule, especially in regards to reactivity and bonding capabilities. Specifically, electron density affects an atom's ability to form bonds, including ionic, covalent, and hydrogen bonds.

Within a molecule, areas of high electron density might attract atoms or ions of opposite charge or repel those with similar charges. For example, negative ions have high electron density due to the extra electrons. In the context of the exercise, the compounds were analyzed based on their varying electron densities to determine which was the strongest Brønsted base. The one with the highest negative charge (in this case, the \([:\mathrm{O} \equiv \mathrm{O}:]^{2-}\) ion) exhibited the highest electron density, hence a greater ability to attract the positively charged protons (H+ ions).

This relationship between electron density and proton affinity is imperative for understanding basicity in chemistry. The compound's electron density facilitates our prediction of its behavior as an acid or base in chemical reactions.

Chemical Bonding

Chemical bonding is the force that holds atoms together in molecules and compounds. There are several types of chemical bonds, including covalent, ionic, metallic, and hydrogen bonds, each with unique properties and behaviors. Covalent bonds, in particular, involve the sharing of electron pairs between atoms. The Lewis structures used in the exercise exemplify such sharing of electrons to achieve stable electron configurations.

In the context of the given exercise, covalent bonds are represented by the triple bonds between the X and Y atoms in the general formula \[[: \mathrm{X} \equiv \mathrm{Y}:]^{n}\]. The strength and nature of these bonds can influence the molecule's reactivity and its subsequent interactions. Lone pairs of electrons, likewise denoted in Lewis structures, play an important role in base strength and in reactions where a molecule acts as a Lewis base by donating a pair of electrons.

Understanding the concepts of electron sharing and distribution in bonds is critical when predicting molecular geometry, polarity, intermolecular forces, and reactivity. Ultimately, the study of chemical bonding not only aids in drawing accurate Lewis structures but also enhances comprehension of the underlying principles governing all chemical compounds and their interactions, such as the strength of a molecule as a Brønsted base.